Kinetics of an Iodine Clock Reaction Lab_ Teacher s Key Purpose: In this lab, you will find the reaction rate, rate law,, and observe the effects of a catalyst for the oxidation of iodide ions by bromate ions in the presence of acid (reaction A). We can measure the rate of a reaction by monitoring the depletion of a reactant s concentration over time or by measuring the formation of a product over time. In this lab, you will use a second reaction that acts as a clock that will provide you a way to measure the time it takes for the reaction to proceed. A: 6I + BrO 3 + 6H + + Starch 3I 2 + Br + 3H 2 O B: I 2 + 2S 2 O 3 2 2I + S 4 O 6 2 The rate of reaction A can be expressed as Δ[BrO 3 ]/Δt. As reaction A produces I 2, it will react with S 2 O 3 2 (reaction B). When S 2 O 3 2 is consumed, excess I 2 will react with starch and the solution will turn blue. You will find the rate by recording the color change time and use stoichiometry to determine the moles of BrO 3 that reacted. By dividing the number of moles of BrO 3 that reacted by the time it took for the solution to turn blue, you will determine the rate of reaction. You will then repeat this experiment varying reactant concentrations and recording the rate. Using your data, you will calculate the reactant order, write the rate law, and lastly you will observe the effects on the reaction rate using a catalyst. Materials: Copper (II) Nitrate Solution,.1M, 5 ml Distilled or DI water, 5 ml Hydrochloric acid solution,.10m, 5 ml Potassium Iodide, KI, 0.010M, 5 ml Potassium bromate, 0.040 M, 5 ml 2% Starch solution, 5 ml Sodium thiosulfate, 0.0010 M, 5 ml Cassette Tape Case Timer (or cell phone timer) Analytical balance Beaker, 10mL or 50mL Beraltype pipet with microtip, 7 Cotton swabs for cleaning well plate Reaction Strips, 12well, 2 Label Tape for pipets Safety: Wear safety goggles and gloves. Keep hair tied back, and wear long sleeve shirts and closedtoed shoes. Wash hands with soap and water before and after doing the lab. Prelab questions: I 2 + 2S 2 O 3 2 2I + S 4 O 6 2 Trial [I 2 ] [S 2 O 2 3 ] Reaction Time 1 0.040 M 0.040 M 300 sec 2 0.080 M 0.040 M 150 sec 3 0.040 M 0.080 M 145 sec 1) In each trial, the blue color appeared after adding 0.0020 M iodide (I ) had been produced. Calculate the reaction rate for each trial by dividing the concentration of iodine formed by the reaction time. Trial 1: rate= 6.70 x 10 6 M s 1 Trial 2: rate= 1.34 x 10 5 M s 1 Trial 3: rate = 1.34 x 10 5 M s 1 2) Find the order of reaction with respect to iodide ions. 1 st order 3) Find the order of reaction with respect to thiosulfate ions. 1 st order 4) Calculate the rate constant for this reaction. k= 4.19 x 10 3 M 1 s 1 5) Write the rate law for this portion of the iodine clock reaction. Can you predict the rate law by looking at just the coefficients? rate =k[i 2 ][S 2 O 2 3 ] You cannot predict the rate law by looking at the 1
coefficients. The exponents are not the same as the coefficients in the balanced equation. Rate laws must be determined experimentally. Procedure: Part 1: You will be using a set number of drops of each reactant. If you know the molarity of the reactant and the mass of 1 drop, you can calculate the volume of the drop (assuming density of the solution is 1.00 g/ml) and thus calculate the moles of each reactant. 1. Obtain a microtip pipet and fill pipet with 3 ml of water (up to 3 ml line). 2. Find the mass of a small beaker using a balance. Record below. 3. Hold pipet vertical and deliver 5 drops of water into beaker and find total mass. Record below. 4. Add five more drops of water into beaker and determine mass. Record below. 5. Add five more drops to the same beaker and find the mass. Record below. Data Table and Calculations Part 1: Trial 1 1. Mass of empty beaker (a): 28.57 g 2. Mass of beaker + 5 drops water (b): 28.65 g 3. Mass of first 5 drops water (b)(a): 0.08 g 4. Average mass of 1 drop of water: 0.016 g Trial 2 1. Mass of beaker plus 10 drops of water (c): 28.73 g 2. Mass of second 5 drops of water (c) (b): 0.08 g 3. Average mass of 1 drop of water: 0.016 g Trial 3 Mass of beaker + 15 drops water (d): 28.82 g Mass of third 5 drops water (d)(c): 0.09 g Average mass of 1 drop of water: 0.018 g Average mass of 1 drop of water (Trials 13): 0.017 g Volume of one drop: (density = 1.00g/mL): 0.017 ml Or 1.7 x 10 5 L Part 2: Determine reaction rate and calculate the rate law. Be careful to be consistent when adding drops of solution to each well. Hold pipets vertically and check to see that there are no air bubbles present. 1) Take 6 microtip pipets and fold a piece of tape around the stem of each pipet. Label each pipet as KI, H 2 O, HCl, Starch, Na 2 S 2 O 3, KBrO 3 2) Fill each pipet with about 2 ml of each of the respective solutions. 3) Place the pipets in an opened cassette case. 4) Obtain two clean, 12well reaction strips. Place them on a clean sheet of white paper and number them from 16. A) Varying the concentration of KI Use the following table as a guide for filling wells 16 in reaction strip 1+2. Mix the solution in each well with a toothpick. Do each experiment twice. Reaction Strip 1 : wells 16 Experiment 1 Experiment 1 Experiment 2 Experiment 2 Experiment 3 Experiment 3 Well # 1 2 3 4 5 6 Drops KI 2 2 4 4 6 6 Drops 4 4 2 2 0 0 Water Drops HCl 2 2 2 2 2 2 Drops Starch 1 1 1 1 1 1 2
Drops Na 2 S 2 0 3 1 1 1 1 1 1 Reaction Strip 2: wells 16 Experiment 1 Experiment 1 Experiment 2 Experiment 2 Experiment 3 Experiment 3 Well # 1 2 3 4 5 6 Drops 2 2 2 2 2 2 KBrO 3 1) Turn the second reaction strip upside down and place it on top of the first strip so that the numbered wells are lined up on top of each other. Surface tension will prevent the liquid from flowing out of the wells. 2) Hold the aligned strips firmly together open end to open end. Shake them downward once with a sharp downward motion. Have one lab partner begin timing immediately when the strips are mixed together. 3) Record the time when the solution in each cell turned blue in the part 2 data table below. 4) When all cells have turned blue, take the temperature of one of the reaction solutions. Record and enter in the part 2 data table below. 5) Rinse contents of the well strips with warm tap water. Use cotton swabs to dry inside of each well. 6) Repeat the entire process (steps 411) for experiments 47. B) Varying the concentration of BrO 3 Use the following table as a guide for filling wells 16 in reaction strip 1+2. Mix the solution in each well with a toothpick. Do each experiment twice. Reaction Strip 1: (experiments 4 and 5) Experiment 4 Experiment 4 Experiment 5 Experiment 5 Well # 1 2 3 4 Drops KI 2 2 2 2 Drops Water 2 2 0 0 Drops HCl 2 2 2 2 Drops Starch Drops Na 2 S 2 0 3 1 1 1 1 1 1 1 1 Reaction Strip 2: wells 16 Experiment 4 Experiment 4 Experiment 5 Experiment 5 Well # 1 2 3 4 Drops 4 4 6 6 KBrO 3 C) Varying the concentration of HCl Use the following table as a guide for filling wells 16 in reaction strip 1+2. Mix the solution in each well with a toothpick. Do each experiment twice. 3
Reaction Strip 1 : (experiments 6 and 7) Experiment 6 Experiment 6 Experiment 7 Experiment 7 Well # 1 2 3 4 Drops KI 2 2 2 2 Drops Water 2 2 0 0 Drops HCl 4 4 6 6 Drops Starch Drops Na 2 S 2 0 3 1 1 1 1 1 1 1 1 Reaction Strip 2: wells 16 Experiment 4 Experiment 4 Experiment 5 Experiment 5 Well # 1 2 3 4 Drops 2 2 2 2 KBrO 3 Part 2 Data Table: Experiment No. Trial 1 Rate (s) Trial 2 Rate (s) Average Rate (s) Temp. C 1 180 175 177.5 23 2 85 83 84 23 3 63 50 56.5 23 4 89 81 85 23 5 59 63 61 23 6 36 44 40 23 7 18 20 19 23 Part 3: Observe the Effect of a Catalyst on the Rate Repeat the procedure given in part 2 for experiment 1 only. This time, however, add 1 drop of.1m Copper (II) nitrate solution and only 3 drops of water to the mixtures. Fill only reaction wells 12. Record the reaction times in the part 3 data table. Part 3 Data Table: Observe the effect of a catalyst on the rate. Reaction Time, Seconds Experiment 1 Catalyzed 168 Experiment 1 Uncatalyzed 54 4
Calculations: 1) Calculate the rate: The rate is the amount of bromate reacted divided by time in seconds. Hint: Calculate the moles of thiosulfate in one drop of 0.0010 M Na 2 S 2 O 3 solution, and use that quantity to calculate the number of moles of BrO 3 that reacted. To find the concentration of BrO 3, divide by the volume of 12 drops (total volume of the solution) and subtract this concentration from the initial concentration. Record in table below. In each reaction there is one drop of 0.0010 M Na 2 S 2 O 3 a) calculate the moles of S 2 O 3 2 in one drop: M=mol/L mol= MxL = (0.0010M) (1.70 x 10 5 L)= 1.70 x 10 8 mol b) When the blue color appears, all of the S 2 O 3 2 has been consumed. So use the number of moles of S 2 O 3 2 that reacted to find the number of moles of I 2 that reacted (or was produced) then use that to find moles of BrO 3 that reacted. 1.7 x 10 8 mol S 2 O 3 2 x 1 I 2 = 8.5 x 10 9 mol of I 2 x 1 mol = 2.8 x 10 9 mol of BrO 3 2 2 S 2 O 3 3 mol I 2 reacted Δ[BrO 3 ] = (2.8 x 10 9 mol of BrO 3 )/ volume of 12 drops = (2.8 x 10 9 mol of BrO 3 )/ 12(1.7 x 10 5 L) = 1.4 x 10 5 M BrO 3 Exp. 1 Exp. 2 Exp. 3 Exp. 4 Exp. 5 Exp. 6 Exp. 7 Reaction Rate, M/s 1.4 x 10 5 M BrO 3 /177.5 = 7.9 x 10 8 M/s 1.4 x 10 5 M BrO 3 /84 = 1.7 x 10 7 M/s 1.4 x 10 5 M BrO 3 / 56.5 = 2.5 x 10 7 M/s 1.4 x 10 5 M BrO 3 / 85 = 1.6 x 10 7 M/s 1.4 x 10 5 M BrO 3 / 61 = 2.3 x 10 7 M/s 1.4 x 10 5 M BrO 3 /40 = 3.5 x 10 7 M/s 1.4 x 10 5 M BrO 3 /19 =7.4 x 10 7 M/s 5
2) Calculate the initial concentration of each reactant and record in the data table below. Hint: Use the M 1 V 1 =M 2 V 2 equation to find the concentration of each reactant (I, BrO 3, and H + ) after it was diluted. Record in table below. Experiment 1: BrO 3 (0.040)(2 drops x 1.75 x 10 5 L)= (M2)(12(1.75 x 10 5 L) => M2= 6.7 x 10 3 M I (0.010)(2 drops x 1.75 x 10 5 L)= (M2)(12(1.75 x 10 5 L) => M2= 1.7 x 10 3 M H + (0.10)(2 drops x 1.75 x 10 5 L)= (M2)(12(1.75 x 10 5 L) => M2= 1.7 x 10 2 M Experiment 2: BrO 3 (0.040)(2 drops x 1.75 x 10 5 L)= (M2)(12(1.75 x 10 5 L) => M2= 6.7 x 10 3 M I (0.010)(4 drops x 1.75 x 10 5 L)= (M2)(12(1.75 x 10 5 L) => M2= 3.4 x 10 3 M H + (0.10)(2 drops x 1.75 x 10 5 L)= (M2)(12(1.75 x 10 5 L) => M2= 1.7 x 10 2 M Experiment 3: BrO 3 (0.040)(2 drops x 1.75 x 10 5 L)= (M2)(12(1.75 x 10 5 L) => M2= 6.7 x 10 3 M I (0.010)(6 drops x 1.75 x 10 5 L)= (M2)(12(1.75 x 10 5 L) => M2= 5.1 x 10 3 M H + (0.10)(2 drops x 1.75 x 10 5 L)= (M2)(12(1.75 x 10 5 L) => M2= 1.7 x 10 2 M Experiment 4: BrO 3 (0.040)(4 drops x 1.75 x 10 5 L)= (M2)(12(1.75 x 10 5 L) => M2= 1.3 x 10 2 M I (0.010)(2 drops x 1.75 x 10 5 L)= (M2)(12(1.75 x 10 5 L) => M2= 1.7 x 10 3 M H + (0.10)(2 drops x 1.75 x 10 5 L)= (M2)(12(1.75 x 10 5 L) => M2= 1.7 x 10 2 M Experiment 5: BrO 3 (0.040)(6 drops x 1.75 x 10 5 L)= (M2)(12(1.75 x 10 5 L) => M2= 2.0 x 10 2 M I (0.010)(2 drops x 1.75 x 10 5 L)= (M2)(12(1.75 x 10 5 L) => M2= 1.7 x 10 3 M H + (0.10)(2 drops x 1.75 x 10 5 L)= (M2)(12(1.75 x 10 5 L) => M2= 1.7 x 10 2 M Experiment 6: BrO 3 (0.040)(2 drops x 1.75 x 10 5 L)= (M2)(12(1.75 x 10 5 L) => M2= 6.7 x 10 3 M I (0.010)(2 drops x 1.75 x 10 5 L)= (M2)(12(1.75 x 10 5 L) => M2= 1.7 x 10 3 M H + (0.10)(2 drops x 1.75 x 10 5 L)= (M2)(12(1.75 x 10 5 L) => M2= 3.4 x 10 2 M Experiment 7: BrO 3 (0.040)(2 drops x 1.75 x 10 5 L)= (M2)(12(1.75 x 10 5 L) => M2= 6.7 x 10 3 M I (0.010)(2 drops x 1.75 x 10 5 L)= (M2)(12(1.75 x 10 5 L) => M2= 1.7 x 10 3 M H + (0.10)(2 drops x 1.75 x 10 5 L)= (M2)(12(1.75 x 10 5 L) => M2= 5.1 x 10 3 M 6
7.9 x 10 8 M/s 7.9 x 10 8 M/s Initial Concentrations, Moles/Liter and Rate of Reaction Experiment [ I ] [BrO 3 ] [H + ] rate 7.9 x 10 8 M/s 7.9 x 1 10 8 M/s 1.7 x 10 3 M 6.7 x 10 3 M 1.7 x 10 2 M 7.9 x 10 8 M/s 2 3.4 x 10 3 M 6.7 x 10 3 M 1.7 x 10 2 M 1.7 x 10 7 M/s 3 5.1 x 10 3 M 6.7 x 10 3 M 1.7 x 10 2 M 2.5 x 10 7 M/s 4 1.7 x 10 3 M 1.3 x 10 2 M 1.7 x 10 2 M 1.6 x 10 7 M/s 5 1.7 x 10 3 M 2.0 x 10 2 M 1.7 x 10 2 M 2.3 x 10 7 M/s 6 1.7 x 10 3 M 6.7 x 10 3 M 3.4 x 10 2 M 3.5 x 10 7 M/s 7 1.7 x 10 3 M 6.7 x 10 3 M 5.1 x 10 3 M 7.4 x 10 7 M/s 3) Calculate the order of each reactant using the data in the table above, and write the rate law for the oxidation of iodide ions by bromate ions in the presence of acid. A) Order of [ I ] = 1.7 x 10 7 M/s = [ 3.4 x 10 3 M] n = 2.1 = 2.10 n n= 1.0 7.9 x 10 8 M/s [1.7 x 10 3 M] n Order of [ I ]= 1.0 B) Order of [BrO 3 ] = 2.3 x 10 7 M/s = [ 2.0 x 10 2 M] n = 1.4 = 1.5 n n= 1.0 1.6 x 10 7 M/s [1.3 x 10 2 M] n Order of [BrO 3 ]= 1.0 C) Order of [H + ] = 7.9x 10 8 M/s = [ 1.7 x 10 2 M] n =.23 =. 50 3.5 x 10 7 M/s [3.4x 10 2 M] n log.23 = n log.50.63 = n (.3) n = 2 Order of [H + ]= 2.0 7
Rate Law = k [ I ] [BrO 3 ] [H + ] 2 5) Calculate the Rate Constant. Don t forget your units! Experiment 1 : 7.9 x 10 8 M/s = k [1.7 x 10 3 M] [6.7 x 10 3 M] [1.7 x 10 2 M] 2 k= 24 M 3 s 1 Experiment 2 : 1.7 x 10 7 M/s = k [3.4 x 10 3 M] [6.7 x 10 3 M] [1.7 x 10 2 M] 2 k= 25 M 3 s 1 Experiment 3 : 2.5 x 10 7 M/s = k [5.1 x 10 3 M ][6.7 x 10 3 M] [1.7 x 10 2 M] 2 k=26 M 3 s 1 Experiment 4 : 1.6 x 10 7 M/s = k [1.7 x 10 3 M] [1.3 x 10 2 M] [1.7 x 10 2 M] 2 k=25 M 3 s 1 Experiment 5 : 2.3 x 10 7 M/s = k[1.7 x 10 3 M] [ 2.0 x 10 2 M] [1.7 x 10 2 M] 2 k=23 M 3 s 1 Experiment 6 : 3.5 x 10 7 M/s = k [1.7 x 10 3 M] [6.7 x 10 3 M ] [3.4 x 10 2 M] 2 k=27 M 3 s 1 Experiment 7: 7.4 x 10 7 M/s = k [1.7 x 10 3 M] [6.7 x 10 3 M] [5.1 x 10 3 M] 2 k=27 M 3 s 1 Experiment Value of k 1 24 M 3 s 2 25 M 3 s 3 26 M 3 s 4 25 M 3 s 1 5 23 M 3 s 1 6 27 M 3 s 1 7 27 M 3 s 1 Average Value of k: 25.2 M 3 s 1 8
Questions: 1) Explain the general procedure you used to find the rate law. In this lab, we used a second equation to help us measure the rate of the first reaction. A: 6I + BrO 3 + 6H + + Starch 3I 2 + Br + 3H 2 O B: I 2 + 2S 2 O 3 2 2I + S 4 O 6 2 The rate of reaction A can be expressed as Δ[BrO 3 ]/Δt. As reaction A produces I 2, it will react with 2 S 2 O 3 (reaction B). When S 2 O 2 3 is consumed, excess I 2 will react with starch and the solution will turn blue. We found the rate by recording the color change time and use stoichiometry to determine the moles of BrO 3 that reacted. By dividing the number of moles of BrO 3 that reacted by the time it took for the solution to turn blue we found the rate for each experiment. We varied the concentration of each reactant so that we could calculate the reactant order. 2) Explain what happened to the activation energy of the reaction when a catalyst is added? Draw a reaction energy diagram without the catalyst. Label both axes. On this graph overlay the line representing the change in activation energy in the presence of a catalyst. Label each line as with catalyst and without catalyst. The catalyst lowered the activation energy and that is why the rate of the reaction decreasesd. 3) Check with your teacher for the known rate law for the oxidation of iodide ions by bromate ions in the presence of acid. What is the overall order of your experimental reaction? What is the published overall order for this reaction? 4 4) Calculate the percent error in the overall order of your reaction. List 4 carefully explained reasons that account for your error. If you think you made a calculation error, you need to redo the work. Answers will vary depending on the students. Causes of error may include adding to much of a given reactant to not adding enough. 9
10